At $25^{\circ}C$,the dissociation constants of acetic acid and $HCN$ are $1.5 \times 10^{-5}$ and $4.5 \times 10^{-10}$ respectively. What is the equilibrium constant for the following reaction? $(CN^- + CH_3COOH \rightleftharpoons HCN + CH_3COO^-)$

The first and second ionization constants of a weak dibasic acid $H_{2}A$ are $8.1 \times 10^{-8}$ and $1.0 \times 10^{-13}$ respectively. $0.1 \text{ mol}$ of $H_{2}A$ was dissolved in $1 \text{ L}$ of $0.1 \text{ M}$ $HCl$ solution. The concentration of $HA^{-}$ in the resultant solution is:

$0.01 \ M \ HA(aq.)$ is $2 \%$ ionized,$[OH^{-}]$ of solution is :-

The first ionization constant of $H_2S$ is $9.1 \times 10^{-8}$. Calculate the concentration of $HS^{-}$ ion in its $0.1 \ M$ solution. How will this concentration be affected if the solution is $0.1 \ M$ in $HCl$ also? If the second dissociation constant of $H_2S$ is $1.2 \times 10^{-13}$,calculate the concentration of $S^{2-}$ under both conditions.

When $300 \, mL$ of $M/3 \, HCl$,$200 \, mL$ of $M/2 \, HNO_3$,and $400 \, mL$ of $M/4 \, NaOH$ are mixed,and the total volume is made up to $1 \, dm^3$,what will be the $pH$ of the resulting solution?

The correct order of $pH$ for $0.1 \ M$ solutions is: